gas

1. 1. The state of matter distinguished from the solid and liquid states by relatively low density and viscosity, relatively great expansion and contraction with changes in pressure and temperature, the ability to diffuse readily, and the spontaneous tendency to become distributed uniformly throughout any container.
   2. A substance in the gaseous state.
2. A gaseous fuel, such as natural gas.
3. Gasoline.
4. The speed control of a gasoline engine. Used with the: Step on the gas.
5. A gaseous asphyxiant, irritant, or poison.
6. A gaseous anesthetic, such as nitrous oxide.
7. 1. Flatulence.
   2. Flatus.
8. Slang. Idle or boastful talk.
9. Slang. Someone or something exceptionally exciting or entertaining: The party was a gas.

1. To treat chemically with gas.
2. To overcome, disable, or kill with poisonous fumes.

1. To give off gas.
2. Slang. To talk excessively.

gas up

1. To supply a vehicle with gas or gasoline: gas up a car; gassed up before the trip.

[Dutch, an occult physical principle supposed to be present in all bodies, alteration of Greek khaos, chaos, empty space, coined by Jan Baptista van Helmont (1577-1644), Flemish chemist.]

For more information on gas, visit Britannica.com.

Concept

The number of elements that appear ordinarily in the form of a gas is relatively small: oxygen, hydrogen, fluorine, and chlorine in the halogen "family"; and a handful of others, most notably the noble gases in Group 8 of the periodic table. Yet many substances can exist in the form of a gas, depending on the relative attraction and motion of molecules in that substance. A simple example, of course, is water, or H₂O, which, though it appears as a liquid at room temperature, begins to vaporize and turn into steam at 212°F (100°C). In general, gases respond more dramatically to changes in pressure and temperature than do most other types of matter, and this allows scientists to predict gas behaviors under certain conditions. These predictions can explain mundane occurrences, such as the fact that an open can of soda will soon lose its fizz, but they also apply to more dramatic, life-and-death situations.

How It Works

Molecular Motion and Phases of Matter

On Earth, three principal phases or states of matter exist: solid, liquid, and gas. The differences between these three are, on the surface at least, easily perceivable. Clearly water is a liquid, just as ice is a solid and steam a vapor or gas. Yet the ways in which various substances convert between phases are often complex, as are the interrelations between these phases. Ultimately, understanding of the phases depends on an awareness of what takes place at the molecular level.

All molecules are in motion, and the rate of that motion determines the attraction between them. The movement of molecules generates kinetic energy, or the energy of movement, which is manifested as thermal energy. In everyday language, thermal energy is what people mean when they say "heat"; but in scientific terms, heat has a different definition.

The force that attracts atoms to atoms, or molecules to molecules, is not the same as gravitational force, which holds the Moon in orbit around Earth, Earth in orbit around the Sun, and so on. By contrast, the force of interatomic and intermolecular attraction is electromagnetic. Just as the north pole of a magnet is attracted to the south pole of another magnet and repelled by that other magnet's north pole, so positive electric charges are attracted to negative charges, and negatives to positives. (In fact, electricity and magnetism are both manifestations of an electromagnetic interaction.)

The electromagnetic attractions between molecules are much more complex than this explanation makes it seem, and they play a highly significant role in chemical bonding. In simple terms, however, one can say that the greater the rate of motion for the molecules in relation to one another, the less the attraction between molecules. In addition, the kinetic energy, and hence the thermal energy, is greater in a substance whose molecules are relatively free to move.

When the molecules in a material move slowly in relation to one another, they exert a strong attraction, and the material is called a solid. Molecules of liquid, by contrast, move at moderate speeds and exert a moderate attraction. A material substance whose molecules move at high speeds, and therefore exert little or no attraction, is known as a gas.

Comparison of Gases to Other Phases of Matter

Water and Air Compared

Gases respond to changes in pressure and temperature in a manner remarkably different from that of solids or liquids. Consider the behavior of liquid water as compared with air—a combination of oxygen (O₂), nitrogen (N₂), and other gases—in response to experiments involving changes in pressure and temperature.

In the first experiment, both samples are subjected to an increase in pressure from 1 atm (that is, normal atmospheric pressure at sea level) to 2 atm. In the second, both experience an increase in temperature from 32°F (0°C) to 212°F (100°C). The differences in the responses of water and air are striking.

A sample of water that experiences an increase in pressure from 1 to 2 atm will decrease in volume by less than 0.01%, while a temperature increase from the freezing point to the boiling point will result in only a 2% increase in volume. For air, however, an equivalent pressure increase will decrease the volume by a whopping 50%, and an equivalent temperature increase results in a volume increase of 37%.

Air and other gases, by definition, have a boiling point below room temperature. If they did not boil and thus become gas well below ordinary temperatures, they would not be described as substances that are in the gaseous state in most circumstances. The boiling point of water, of course, is higher than room temperature, and that of solids is much higher.

The Arrangement of Particles

Solids possess a definite volume and a definite shape, and are relatively noncompressible: for instance, if one applies extreme pressure to a steel plate, it will bend, but not much. Liquids have a definite volume, but no definite shape, and tend to be noncompressible. Gases, on the other hand, possess no definite volume or shape, and are highly compressible.

At the molecular level, particles of solids tend to be definite in their arrangement and close in proximity—indeed, part of what makes a solid "solid," in the everyday meaning of that term, is the fact that its constituent parts are basically immovable. Liquid molecules, too, are close in proximity, though random in arrangement. Gas molecules are random in arrangement, but tend to be more widely spaced than liquid molecules.

Pressure

There are a number of statements, collectively known as the "gas laws," that describe and predict the behavior of gases in response to changes in temperature, pressure, and volume. Temperature and volume are discussed elsewhere in this book. However, the subject of pressure requires some attention before we can continue with a discussion of the gas laws.

When a force is applied perpendicular to a surface area, it exerts pressure on that surface. Hence the formula for pressure is p = F/A, where p is pressure, F force, and A the area over which the force is applied. The greater the force, and the smaller the area of application, the greater the pressure; conversely, an increase in area—even without a reduction in force—reduces the overall pressure.

Pressure is measured by a number of units in the English and SI systems. Because p = F/A, all units of pressure represent some ratio of force to surface area.

Units of Pressure

The principal SI unit of pressure is called a pascal (Pa), or 1 N/m². It is named for French mathematician and physicist Blaise Pascal (1623-1662), who is credited with Pascal's principle. The latter holds that the external pressure applied on a fluid—which, in the physical sciences, can mean either a gas or a liquid—is transmitted uniformly throughout the entire body of that fluid.

A newton (N), the SI unit of force, is equal to the force required to accelerate 1 kg of mass at a rate of 1 m/sec². Thus a Pascal is the pressure of 1 newton over a surface area of 1 m². In the English or British system, pressure is measured in terms of pounds per square inch, abbreviated as lbs./in². This is equal to 6.89 · 10³ Pa, or 6,890 Pa.

Another important measure of pressure is the atmosphere (atm), which is the average pressure exerted by air at sea level. In English units, this is equal to 14.7 lb/in², and in SI units, to 1.013 · 10⁵ Pa.

There are two other specialized units of pressure measurement in the SI system: the bar, equal to 10⁵ Pa, and the torr, equal to 133 Pa. Meteorologists, scientists who study weather patterns, use the millibar (mb), which, as its name implies, is equal to 0.001 bars. At sea level, atmospheric pressure is approximately 1,013 mb.

The torr, also known as the millimeter of mercury (mm Hg), is the amount of pressure required to raise a column of mercury (chemical symbol Hg) by 1 mm. It is named for Italian physicist Evangelista Torricelli (1608-1647), who invented the barometer, an instrument for measuring atmospheric pressure.

The Barometer

The barometer constructed by Torricelli in 1643 consisted of a long glass tube filled with mercury. The tube was open at one end, and turned upside down into a dish containing more mercury: the open end was submerged in mercury, while the closed end at the top constituted a vacuum—that is, an area devoid of matter, including air.

The pressure of the surrounding air pushed down on the surface of the mercury in the bowl, while the vacuum at the top of the tube provided an area of virtually no pressure into which the mercury could rise. Thus the height to which the mercury rose in the glass tube represented normal air pressure (that is, 1 atm.) Torricelli discovered that at standard atmospheric pressure, the column of mercury rose to 760 mm (29.92 in).

The value of 1 atm was thus established asequal to the pressure exerted on a column ofmercury 760 mm high at a temperature of 0°C(32°F). In time, Torricelli's invention became afixture both of scientific laboratories and ofhouseholds. Since changes in atmospheric pressure have an effect on weather patterns, manyhome indoor-outdoor thermometers today alsoinclude a barometer.

Real-Life Applications

Introduction to the Gas Laws

English chemist Robert Boyle (1627-1691), who made a number of important contributions to chemistry—including his definition and identification of elements—seems to have been influenced by Torricelli. If so, this is an interesting example of ideas passing from one great thinker to another: Torricelli, a student of Galileo Galilei (1564-1642), was no doubt influenced by Galileo's thermoscope.

Like Torricelli, Boyle conducted tests involving the introduction of mercury to a tube closed at the other end. The tube Boyle used was shaped like the letter J, and it was so long that he had to use the multi-story foyer of his house as a laboratory. At the tip of the curved bottom was an area of trapped gas, and into the top of the tube, Boyle introduced increasing quantities of mercury. He found that the greater the volume of mercury, the greater the pressure on the gas, and the less the volume of gas at the end of the tube. As a result, he formulated the gas law associated with his name.

The gas laws are not a set of government regulations concerning use of heating fuel; rather, they are a series of statements concerning the behavior of gases in response to changes in temperature, pressure, and volume. These were derived, beginning with Boyle's law, during the seventeenth, eighteenth, and nineteenth centuries by scientists whose work is commemorated through the association of their names with the laws they discovered. In addition to Boyle, these men include fellow English chemists John Dalton (1766-1844) and William Henry (1774-1836); French physicists and chemists J. A. C. Charles (1746-1823) and Joseph Gay-Lussac (1778-1850); and Italian physicist Amedeo Avogadro (1776-1856).

There is a close relationship between Boyle's, Charles's, and Gay-Lussac's laws. All of these treat one of three parameters—temperature, pressure, or volume—as fixed quantities in order to explain the relationship between the other two variables. Avogadro's law treats two of the parameters as fixed, thereby establishing a relationship between volume and the number of molecules in a gas. The ideal gas law sums up these four laws, and the kinetic theory of gases constitutes an attempt to predict the behavior of gases based on these laws. Finally, Dalton's and Henry's laws both relate to partial pressure of gases.

Boyle's, Charles's, and Gay-Lussac's Laws

Boyle's and Charles's Laws

Boyle's law holds that in isothermal conditions (that is, a situation in which temperature is kept constant), an inverse relationship exists between the volume and pressure of a gas. (An inverse relationship is a situation involving two variables, in which one of the two increases in direct proportion to the decrease in the other.) In this case, the greater the pressure, the less the volume and vice versa. Therefore, the product of the volume multiplied by the pressure remains constant in all circumstances.

Charles's law also yields a constant, but in this case the temperature and volume are allowed to vary under isobarometric conditions—that is, a situation in which the pressure remains the same. As gas heats up, its volume increases, and when it cools down, its volume reduces accordingly. Hence, Charles established that the ratio of temperature to volume is constant.

Absolute Temperature

In about 1787, Charles made an interesting discovery: that at 0°C (32°F), the volume of gas at constant pressure drops by 1/273 for every Celsius degree drop in temperature. This seemed to suggest that the gas would simply disappear if cooled to −273°C (−459.4°F), which, of course, made no sense. In any case, the gas would most likely become first a liquid, and then a solid, long before it reached that temperature.

The man who solved the quandary raised by Charles's discovery was born a year after Charles died. He was William Thomson, Lord Kelvin (1824-1907); in 1848, he put forward the suggestion that it was molecular translational energy—the energy generated by molecules in motion—and not volume, that would become zero at −273°C. He went on to establish what came to be known as the Kelvin scale of absolute temperature.

Sometimes known as the absolute temperature scale, the Kelvin scale is based not on the freezing point of water, but on absolute zero—the temperature at which molecular motion comes to a virtual stop. This is −273.15°C (−459.67°F). In the Kelvin scale, which uses neither the term nor the symbol for "degree," absolute zero is designated as 0K.

Scientists prefer the Kelvin scale to the Celsius, and certainly to the Fahrenheit, scales. If the Kelvin temperature of an object is doubled, its average molecular translational energy has doubled as well. The same cannot be said if the temperature were doubled from, say, 10°C to 20°C, or from 40°F to 80°F, since neither the Celsius nor the Fahrenheit scale is based on absolute zero.

Gay-Lussac's Law

From Boyle's and Charles's law, a pattern should be emerging: both treat one parameter (temperature in Boyle's, pressure in Charles's) as unvarying, while two other factors are treated as variables. Both, in turn, yield relationships between the two variables: in Boyle's law, pressure and volume are inversely related, whereas in Charles's law, temperature and volume are directly related.

In Gay-Lussac's law, a third parameter, volume, is treated as a constant, and the result is a constant ratio between the variables of pressure and temperature. According to Gay-Lussac's law, the pressure of a gas is directly related to its absolute temperature.

Avogadro's Law

Gay-Lussac also discovered that the ratio in which gases combine to form compounds can be expressed in whole numbers: for instance, water is composed of one part oxygen and two parts hydrogen. In the language of modern chemistry, this is expressed as a relationship between molecules and atoms: one molecule of water contains one oxygen atom and two hydrogen atoms.

In the early nineteenth century, however, scientists had yet to recognize a meaningful distinction between atoms and molecules, and Avogadro was the first to achieve an understanding of the difference. Intrigued by the whole-number relationship discovered by Gay-Lussac, Avogadro reasoned that one liter of any gas must contain the same number of particles as a liter of another gas. He further maintained that gas consists of particles—which he called molecules—that in turn consist of one or more smaller particles.

In order to discuss the behavior of molecules, Avogadro suggested the use of a large quantity as a basic unit, since molecules themselves are very small. Avogadro himself did not calculate the number of molecules that should be used for these comparisons, but when that number was later calculated, it received the name "Avogadro's number" in honor of the man who introduced the idea of the molecule. Equal to 6.022137 · 10²³, Avogadro's number designates the quantity of atoms or molecules (depending on whether the substance in question is an element or a compound) in a mole.

Today the mole (abbreviated mol), the SI unit for "amount of substance," is defined precisely as the number of carbon atoms in 12.01 g of carbon. The term "mole" can be used in the same way we use the word "dozen." Just as "a dozen" can refer to twelve cakes or twelve chickens, so "mole" always describes the same number of molecules. The ratio of mass between a mole of Element A and Element B, or Compound A and Compound B, is the same as the ratio between the mass of Atom A and Atom B, or Molecule A and Molecule B. Avogadro's law describes the connection between gas volume and number of moles. According to Avogadro's law, if the volume of gas is increased under isothermal and isobarometric conditions, the number of moles also increases. The ratio between volume and number of moles is therefore a constant.

The Ideal Gas Law

Once again, it is easy to see how Avogadro's law can be related to the laws discussed earlier. Like the other three, this one involves the parameters of temperature, pressure, and volume, but it also introduces a fourth—quantity of molecules (that is, number of moles). In fact, all the laws so far described are brought together in what is known as the ideal gas law, sometimes called the combined gas law.

The ideal gas law can be stated as a formula, pV = nRT, where p stands for pressure, V for volume, n for number of moles, and T for temperature. R is known as the universal gas constant, a figure equal to 0.0821 atm · liter/mole · K. (Like most figures in chemistry, this one is best expressed in metric rather than English units.)

Given the equation pV = nRT and the fact that R is a constant, it is possible to find the value of any one variable—pressure, volume, number of moles, or temperature—as long as one knows the value of the other three. The ideal gas law also makes it possible to discern certain relationships: thus, if a gas is in a relatively cool state, the product of its pressure and volume is proportionately low; and if heated, its pressure and volume product increases correspondingly.

The Kinetic Theory of Gases

From the preceding gas laws, a set of propositions known collectively as the kinetic theory of gases has been derived. Collectively, these put forth the proposition that a gas consists of numerous molecules, relatively far apart in space, which interact by colliding. These collisions are responsible for the production of thermal energy, because when the velocity of the molecules increases—as it does after collision—the temperature increases as well.

There are five basic postulates to the kinetic theory of gases:

• 1. Gases consist of tiny molecular or atomic particles.
• 2. The proportion between the size of these particles and the distances between them is so small that the individual particles can be assumed to have negligible volume.
• 3. These particles experience continual random motion. When placed in a container, their collisions with the walls of the container constitute the pressure exerted by the gas.
• 4. The particles neither attract nor repel one another.
• 5. The average kinetic energy of the particles in a gas is directly related to absolute temperature.

These observations may appear to resemble statements made earlier concerning the differences between gases, liquids, and solids in terms of molecular behavior. If so, that is no accident: the kinetic theory constitutes a generally accepted explanation for the reasons why gases behave as they do. Kinetic theories do not work as well for explaining the behaviors of solids and liquids; nonetheless, they do go a long way toward identifying the molecular properties inherent in the various phases of matter.

Laws of Partial Pressure

In addition to all the gas laws so far discussed, two laws address the subject of partial pressure. When two or more gases are present in a container, partial pressure is the pressure that one of them exerts if it alone is in the container.

Dalton's law of partial pressure states that the total pressure of a gas is equal to the sum of its partial pressures. As noted earlier, air is composed mostly of nitrogen and oxygen. Along with these are small components, carbon dioxide, and gases collectively known as the rare or noble gases: argon, helium, krypton, neon, radon, and xenon. Hence, the total pressure of a given quantity of air is equal to the sum of the pressures exerted by each of these gases.

Henry's law states that the amount of gas dissolved in a liquid is directly proportional to the partial pressure of the gas above the surface of the solution. This applies only to gases such as oxygen and hydrogen that do not react chemically to liquids. On the other hand, hydrochloric acid will ionize when introduced to water: one or more of its electrons will be removed, and its atoms will convert to ions, which are either positive or negative in charge.

Applications of Dalton's and Henry's Laws

Partial Pressure: a Matter of Life and Possible Death for Scuba Divers

The gas laws are not just a series of abstract statements. Certainly, they do concern the behavior of ideal as opposed to real gases. Like all scientific models, they remove from the equation all outside factors, and treat specific properties in isolation. Yet, the behaviors of the ideal gases described in the gas laws provide a key to understanding the activities of real gases in the real world. For instance, the concept of partial pressure helps scuba divers avoid a possibly fatal sickness.

Imagine what would happen if a substance were to bubble out of one's blood like carbon dioxide bubbling out of a soda can, as described below. This is exactly what can happen to an undersea diver who returns to the surface too quickly: nitrogen rises up within the body, producing decompression sickness—known colloquially as "the bends." This condition may manifest as itching and other skin problems, joint pain, choking, blindness, seizures, unconsciousness, permanent neurological defects such as paraplegia, and possibly even death.

If a scuba diver descending to a depth of 150 ft (45.72 m) or more were to use ordinary air in his or her tanks, the results would be disastrous. The high pressure exerted by the water at such depths creates a high pressure on the air in the tank, meaning a high partial pressure on the nitrogen component in the air. The result would be a high concentration of nitrogen in the blood, and hence the bends.

Instead, divers use a mixture of helium and oxygen. Helium gas does not dissolve well in blood, and thus it is safer for a diver to inhale this oxygen-helium mixture. At the same time, the oxygen exerts the same pressure that it would normally—in other words, it operates in accordance with Dalton's observations concerning partial pressure.

Opening a Soda Can

Inside a can or bottle of carbonated soda is carbon dioxide gas (CO₂), most of which is dissolved in the drink itself. But some of it is in the space (sometimes referred to as "head space") that makes up the difference between the volume of the soft drink and the volume of the container.

At the bottling plant, the soda manufacturer adds high-pressure carbon dioxide (CO₂) to the head space in order to ensure that more CO₂ will be absorbed into the soda itself. This is in accordance with Henry's law: the amount of gas (in this case CO₂) dissolved in the liquid (soda) is directly proportional to the partial pressure of the gas above the surface of the solution—that is, the CO₂ in the head space. The higher the pressure of the CO₂ in the head space, the greater the amount of CO₂ in the drink itself; and the greater the CO₂ in the drink, the greater the "fizz" of the soda.

Once the container is opened, the pressure in the head space drops dramatically. Once again, Henry's law indicates that this drop in pressure will be reflected by a corresponding drop in the amount of CO₂ dissolved in the soda. Over a period of time, the soda will release that gas, and eventually, it will go "flat."

Fire Extinguishers

A fire extinguisher consists of a long cylinder with an operating lever at the top. Inside the cylinder is a tube of carbon dioxide surrounded by a quantity of water, which creates pressure around the CO₂ tube. A siphon tube runs vertically along the length of the extinguisher, with one opening in the water near the bottom. The other end opens in a chamber containing a spring mechanism attached to a release valve in the CO₂ tube.

The water and the CO₂ do not fill the entire cylinder: as with the soda can, there is "head space," an area filled with air. When the operating lever is depressed, it activates the spring mechanism, which pierces the release valve at the top of the CO₂ tube. When the valve opens, the CO₂ spills out in the "head space," exerting pressure on the water. This high-pressure mixture of water and carbon dioxide goes rushing out of the siphon tube, which was opened when the release valve was depressed. All of this happens, of course, in a fraction of a second—plenty of time to put out the fire.

Aerosol Cans

Aerosol cans are similar in structure to fire extinguishers, though with one important difference. As with the fire extinguisher, an aerosol can includes a nozzle that depresses a spring mechanism, which in turn allows fluid to escape through a tube. But instead of a gas cartridge surrounded by water, most of the can's interior is made up of the product (for instance, deodorant), mixed with a liquid propellant.

The "head space" of the aerosol can is filled with highly pressurized propellant in gas form, and, in accordance with Henry's law, a corresponding proportion of this propellant is dissolved in the product itself. When the nozzle is depressed, the pressure of the propellant forces the product out through the nozzle.

A propellant, as its name implies, propels the product itself through the spray nozzle when the nozzle is depressed. In the past, chlorofluorocarbons (CFCs)—manufactured compounds containing carbon, chlorine, and fluorine atoms—were the most widely used form of propellant. Concerns over the harmful effects of CFCs on the environment, however, has led to the development of alternative propellants, most notably hydrochlorofluorocarbons (HCFCs), CFC-like compounds that also contain hydrogen atoms.

Applications of Boyle's, Charles's, and Gay-Lussac's Laws

When the Temperature Changes

A number of interesting results occur when gases experience a change in temperature, some of them unfortunate and some potentially lethal. In these instances, it is possible to see the gas laws—particularly Boyle's and Charles's—at work.

There are numerous examples of the disastrous effects that result from an increase in the temperature of combustible gases, including natural gas and petroleum-based products. In addition, the pressure on the gases in aerosol cans makes the cans highly explosive—so much so that discarded cans at a city dump may explode on a hot summer day. Yet, there are other instances when heating a gas can produce positive effects.

A hot-air balloon, for instance, floats because the air inside it is not as dense than the air outside. According to Charles's law, heating a gas will increase its volume, and since gas molecules exert little attraction toward one another, they tend to "spread out" even further with an increase of volume. This, in turn, creates a significant difference in density between the air in the balloon and the air outside, and as a result, the balloon floats.

Although heating a gas can be beneficial, cooling a gas is not always a wise idea. If someone were to put a bag of potato chips into a freezer, thinking this would preserve their flavor, he would be in for a disappointment. Much of what maintains the flavor of the chips is the pressurization of the bag, which ensures a consistent internal environment so that preservative chemicals, added during the manufacture of the chips, can keep them fresh. Placing the bag in the freezer causes a reduction in pressure, as per Gay-Lussac's law, and the bag ends up a limp version of its former self.

Propane tanks and tires offer an example of the pitfalls that may occur by either allowing a gas to heat up or cool down by too much. Because most propane tanks are made according to strict regulations, they are generally safe, but it is not entirely inconceivable that the extreme heat of a summer day could cause a defective tank to burst. An increase in temperature leads to an increase in pressure, in accordance with Gay-Lussac's law, and could lead to an explosion.

Because of the connection between heat and pressure, propane trucks on the highways during the summer are subjected to weight tests to ensure that they are not carrying too much gas. On the other hand, a drastic reduction in temperature could result in a loss in gas pressure. If a propane tank from Florida were transported by truck during the winter to northern Canada, the pressure is dramatically reduced by the time it reaches its destination.

The Internal-Combustion Engine

In operating a car, we experience two applications of the gas laws. One of these is what makes the car run: the combustion of gases in the engine, which illustrates the interrelation of volume, pressure, and temperature expressed in the laws attributed to Boyle, Charles, and Gay-Lussac. The other is, fortunately, a less frequent phenomenon—but it can and does save lives. This is the operation of an airbag, which depends, in part, on the behaviors explained in Charles's law.

When the driver of a modern, fuel-injection automobile pushes down on the accelerator, this activates a throttle valve that sprays droplets of gasoline mixed with air into the engine. The mixture goes into the cylinder, where the piston moves up, compressing the gas and air. While the mixture is still at a high pressure, the electric spark plug produces a flash that ignites the gasoline-air mixture. The heat from this controlled explosion increases the volume of air, which forces the piston down into the cylinder. This opens an outlet valve, causing the piston to rise and release exhaust gases.

As the piston moves back down again, an inlet valve opens, bringing another burst of gasoline-air mixture into the chamber. The piston, whose downward stroke closed the inlet valve, now shoots back up, compressing the gas and air to repeat the cycle. The reactions of the gasoline and air to changes in pressure, temperature, and volume are what move the piston, which turns a crankshaft that causes the wheels to rotate.

The Airbag

So much for moving—what about stopping? Most modern cars are equipped with an airbag, which reacts to sudden impact by inflating. This protects the driver and front-seat passenger, who, even if they are wearing seatbelts, may otherwise be thrown against the steering wheel or dashboard.

In order to perform its function properly, the airbag must deploy within 40 milliseconds (0.04 seconds) of impact. Not only that, but it has to begin deflating before the body hits it. If a person's body, moving forward at speeds typical in an automobile accident, were to smash against a fully inflated airbag, it would feel like hitting concrete—with all the expected results.

The airbag's sensor contains a steel ball attached to a permanent magnet or a stiff spring. The spring or magnet holds the ball in place through minor mishaps when an airbag is not warranted—for instance, if a car were simply to be "tapped" by another in a parking lot. But in a case of sudden deceleration, the magnet or spring releases the ball, sending it down a smooth bore. The ball flips a switch, turning on an electrical circuit. This in turn ignites a pellet of sodium azide, which fills the bag with nitrogen gas.

At this point, the highly pressurized nitrogen gas molecules begin escaping through vents. Thus, as the driver's or rider's body hits the airbag, the deflation of the bag is moving it in the same direction that the body is moving—only much, much more slowly. Two seconds after impact, which is an eternity in terms of the processes involved, the pressure inside the bag has returned to 1 atm.

The chemistry of the airbag is particularly interesting. The bag releases inert, or non-reactive, nitrogen gas, which poses no hazard to human life; yet one of the chemical ingredients in the airbag is so lethal that some environmentalist groups have begun to raise concerns over its presence in airbags. This is sodium azide (NaN₃), one of three compounds—along with potassium nitrate (KNO₃) and silicon dioxide (SiO₂)—present in an airbag prior to inflation.

The sodium azide and potassium nitrate react to one another, producing a burst of hot nitrogen gas in two back-to-back reactions. In the fractions of a second during which this occurs, the airbag becomes like a solid-rocket booster, experiencing a relatively slow detonation known as "deflagration."

The first reaction releases nitrogen gas, which fills the bag, while the second reaction leaves behind the by-products potassium oxide (K₂O) and sodium oxide (Na₂O). These combine with the silicon dioxide to produce a safe, stable compound known as alkaline silicate. The latter, similar to the sand used for making glass, is all that remains in the airbag after the nitrogen gas has escaped.

Where to Learn More

"Atmospheric Pressure: The Force Exerted by the Weight ofAir" (Web site). <http://kids.earth.nasa.gov/archive/air_pressure/> (April 7, 2001).

"Chemical Sciences Structure: Structure of Matter: Nature of Gases" University of Alberta Chemistry Department (Web site). <http://www.chem.ualberta.ca/~plambeck/che/struct/s0307.htm> (May 12, 2001).

"Chemistry Units: Gas Laws."<http://bio.bio.rpi.edu/MS99/ausemaW/chem/gases.html> (February 21, 2001).

"Homework: Science: Chemistry: Gases" Channelone.com (Web site). <http://www.channelone.com/fasttrack/science/chemistry/gases.html> (May 12, 2001).

"Kinetic Theory of Gases: A Brief Review" University ofVirginia Department of Physics (Web site). <http://www.phys.virginia.edu/classes/252/kinetic_theory.html> (April 15, 2001).

Laws of Gases. New York: Arno Press, 1981.

Macaulay, David. The New Way Things Work. Boston: Houghton Mifflin, 1998.

Mebane, Robert C. and Thomas R. Rybolt. Air and OtherGases. Illustrations by Anni Matsick. New York: Twenty-First Century Books, 1995.

"Tutorials—6." Chemistrycoach.com (Web site). <http://www.chemistrycoach.com/tutorials-6.htm> (February 21, 2001).

Zumdahl, Steven S. Introductory Chemistry: A Foundation, 4th ed. Boston: Houghton Mifflin, 2000.

A phase of matter characterized by relatively low density, high fluidity, and lack of rigidity. A gas expands readily to fill any containing vessel. Usually a small change of pressure or temperature produces a large change in the volume of the gas. The equation of state describes the relation between the pressure, volume, and temperature of the gas. In contrast to a crystal, the molecules in a gas have no long-range order.

At sufficiently high temperatures and sufficiently low pressures, all substances obey the ideal-gas, or perfect-gas, equation of state below, where p is the pressure, T is the absolute temperature, V is the molar volume, and R is the gas constant. Absolute temperature T is expressed on the Kelvin scale. The gas constant is 8.314 joules/(mole K). The molar volume is the molecular weight divided by the gas density.

At lower temperatures and higher pressures, the equation of state of a real gas deviates from that of a perfect gas. Various empirical relations have been proposed to explain the behavior of real gases.

noun

1. Incessant and usually inconsequential talk: babble, blab, blabber, chat, chatter, chitchat, jabber, palaver, prate, prattle, small talk. Slang gab, yak. See words.
2. Something or someone uproariously funny or absurd: absurdity. Informal hoot, joke, laugh, scream. Slang howl, panic, riot. Idioms: a laugh a minute. See laughter.

verb

To talk volubly, persistently, and usually inconsequentially: babble, blabber, chatter, chitchat, clack, jabber, palaver, prate, prattle, rattle (on), run on. Informal go on, spiel. Slang gab, jaw, yak. Idioms: run off at the mouth, shoot thebreezebull. See words.

Idioms beginning with gas:
gasket
gas up

In addition to the idiom beginning with gas, also see cook with gas; run out of steam (gas).

Definition: something not liquid or solid
Antonyms: liquid, solid

[as in ‘gas chamber’]

1. interj. A term of disgust and hatred, implying that gas should be dispensed in generous quantities, thereby exterminating the source of irritation. “Some loser just reloaded the system for no reason! Gas!”

2. interj. A suggestion that someone or something ought to be flushed out of mercy. “The system's getting wedged every few minutes. Gas!”

3. vt. To flush (sense 1). “You should gas that old crufty software.”

4. [IBM] n. Dead space in nonsequentially organized files that was occupied by data that has since been deleted; the compression operation that removes it is called degassing (by analogy, perhaps, with the use of the same term in vacuum technology).

5. [IBM] n. Empty space on a disk that has been clandestinely allocated against future need.

A fluid with no definite volume or shape whose molecules are practically unrestricted by cohesive forces.

Definition

Gas, or flatus, is produced when naturally occurring bacteria in the gastrointestinal tract begin to break down, or digest, food. When an excess of air builds up in the tract from swallowing air or a disorder that prevents digestion, it is released as gas. Gastrointestinal gases include methane, carbon dioxide, nitrogen, and hydrogen.

Description

Gas production is an essential, normal function of the gastrointesinal tract, and most healthy individuals pass up to 1,200 cc (over 40 oz) of gas each day. However, when gas causes excessive pain and cramping (colic) then evaluation and treatment are appropriate.

Causes & Symptoms

Gastrointestinal gas production can be increased by certain foods, illnesses, and some medications. Common causes of excessive gas include:

• Gas-producing foods. Onions, beans, the cabbage family, and other fibrous foods can cause excessive gas or intestinal spasms in some individuals.
• Gastrointestinal diseases and disorders. Increased flatulence is a defining symptom of irritable bowel syndrome, diverticulitis, lactose intolerance, malabsorption problems, dysbiosis (digestive problems), and other gastrointestinal disorders.
• Air swallowing. Swallowing too much air while eating or chewing gum can introduce extra gas to the gastrointestinal tract.
• Medications. Certain prescription and over-the-counter medications may cause gas as a side-effect.
• Stress and food allergies can also cause gas.

Symptoms of excessive gas production include:

• flatulence
• belching, or burping
• abdominal cramping, or colic
• abdominal pain

Diagnosis

A thorough medical and dietary history and physical examination performed by a healthcare professional can usually identify the cause of gas pains resulting from changes to diet or medication. Gas problems triggered by gastrointestinal disease may be harder to diagnose, and will typically require additional medical testing such

as colonoscopy, barium enema, or an upper and/or lower gastrointestinal (GI) series.

Treatment

For excessive gas caused by a particular food or beverage, adjustments to diet can relieve most symptoms. Gas caused by air swallowing can be alleviated by eating more slowly and avoiding gum chewing.

An herbalist or naturopathic healthcare professional may recommend a preparation of a carminative (gas reducing) herb such as valerian (Valeriana officinalis), or peppermint (Mentha piperita), which may be helpful in eliminating discomfort and gas-related bloating.

Homeopathic remedies for excessive intestinal gas include Carbo vegetabilis, Nux vomica, and Chamomilla. The prescription of a specific homeopathic remedy will depend on an individual's overall symptom picture, mood, and temperament, and should only be prescribed by a qualified homeopathic physician.

Hydrotherapy, acupressure, acupuncture, yoga, reflexology, and mild exercise can also help to relieve the pain and discomfort of excessive gas.

Allopathic Treatment

Over-the-counter preparations of the enzyme alpha-D-galactosidase (Beano) can alleviate gas symptoms caused by ingestion of certain foods in some individuals. These preparations are typically available in liquid or tablet form. Other non-prescription medications such as Gas-X, Phazyme, and Mylanta contain the ingredient simethicone, which can reduce gas bubbles within the gastrointestinal tract.

Expected Results

Mild excess gas is typically easy to treat, especially that triggered by dietary causes. Gas caused by gastrointestinal disease may be more difficult to manage, and successful treatment depends on the type and severity of the disorder.

Prevention

Avoiding fermented foods, drastic increases in fiber intake, and excessive air intake can prevent gas in some individuals. Lactose intolerant individuals should avoid dairy products.

Resources

Books

Hoffman, David. The Complete Illustrated Herbal. New York: Barnes & Noble Books, 1999.

Periodicals

Wu, Olivia. "Miss the Bloat: How to Avoid Bloating." Vegetarian Times (June 2000): 80.

Organizations

The National Institute of Diabetes & Digestive & Kidney Diseases (NIDDK). Office of Communications and Public Liaison. NIDDK, National Institutes of Health, 31 Center Drive, MSC 2560, Bethesda, MD 20892-2560. http://www.niddk.nih.gov/index.htm.

[Article by: Paula Ford-Martin]

n. pl. gases or gasses 1. gas or vapor used as a poisonous agent to kill or disable an enemy in warfare.

2. informal short for gasoline.

1. attack with or expose to poisonous gas.

2. kill by exposure to poisonous gas.

See the Introduction, Abbreviations and Pronunciation for further details.

Various legal issues arise concerning the use and distribution of gas.

Supply

A municipal corporation does not have the duty to supply gas to its population. In the event that a city assumes the performance of such function, it is acting merely as a business corporation.

The charter of a gas company is a franchise granted by the state. The manufacture of distribution of gas for light, fuel, or power is a business of a public character, and, therefore, a gas company is ordinarily considered to be a public or quasi-public corporation or a business affected with a public interest. A state may regulate gas companies for the protection of the public and may delegate its regulatory powers to municipal corporations in which gas companies operate. In a number of states, gas companies are subject to a public service commission or other such agency. The jurisdiction of the commission ordinarily includes the power to establish rates and to set forth rules and regulations affecting the service, operation, management, and conduct of the business.

Consumer Supply

Upon obtaining a franchise to supply gas to a particular geographic area, a gas company is bound to fulfill its obligation; it cannot withdraw its service from an area merely because it is dissatisfied with the rates permitted there. Once the franchise of a company has expired, it may withdraw the service. A court may, in certain instances, enjoin the discontinuance of service for a reasonable period — to circumvent undue hardship and inconvenience to the residents of the area.

A gas company has the duty to serve all those who are within the franchise area who desire service and subscribe to the reasonable rules that it may set forth. A municipality or corporation supplying gas may make reasonable rules and regulations to secure the payment of bills, such as eliminating service to the consumer. If there is a genuine controversy about the amount owed, a company is not permitted to discontinue service. A gas company may not require the owner or occupant of a building to pay overdue and unpaid bills by a former owner or occupant before it continues service to the building. Some statutes require that gas companies install a meter on the premises, in order to register the consumption of gas by each customer; and where a customer tampers with the meter and uses a significant amount of unmetered gas, the company can discontinue service and refuse to restore it until the customer pays the amount due for the unmetered gas taken.

A gas company that wrongfully refuses to supply a customer with gas is liable for damages. There are also statutory penalties in some states for such wrongful refusal.

Injuries

A gas company is under the obligation to exercise ordinary care in the construction of its works and the conduct of its business in order to protect life and property.

Gas has a highly dangerous and volatile character and tends to escape. A gas company must, therefore, exercise care to avoid harm to others and is liable for its negligence that results in injury to others by reason of the escape or explosion of gas. It must exercise reasonable care in the inspection of its pipes to ensure that leaks may be discovered promptly; and if leaks or defects in the pipes of the company occur due to faulty construction or maintenance, the company is liable for resulting injuries, even though it did not know about the leak.

In the event that the company has taken due care in the inspection of its pipes and a defect or a break occurs through natural causes or by the act of a third person, the gas company must be given notice of the defect and reasonable time to repair it before liability accrues. A gas company subject to notice that gas is escaping is under an obligation to shut off the gas supply until the necessary repairs have been made.

A gas company has a property right in the mains and pipes and other appliances, and where there is unauthorized interference with, or damage to, this property, the company is entitled to recover damages and an injunction if the circumstances so warrant.

Rates

A gas company has a legal obligation to charge reasonable rates. One of the main purposes of the regulation of gas companies is to prescribe fair and reasonable rates for the selling of gas to the public. Rate increases are permitted only following an impartial and complete investigation — with the object of doing justice to the gas company as well as the public. Relief can be sought in the courts if gas rates are unreasonable — to determine whether the rate making body acted beyond the scope of its power or against the weight of the evidence. The courts, however, cannot decide what rates are reasonable, nor can they put those rates into effect.

See: public utilities.

In physics, one of the phases of matter. The atoms or molecules in gases are more widely spaced than in solids or liquids and suffer only occasional collisions with one another.

Any elastic aeriform fluid in which the molecules are widely separated from each other and so have free paths.

• alveolar g. — the gas in the alveoli of the lungs, where gaseous exchange with the capillary blood takes place. See also oxygen, carbon dioxide.
• blood g. — see blood gas analysis.
• g. bubble disease — a disease of fish in tanks in which the water is supersaturated with oxygen or nitrogen. Gas embolism develops in the gills. Air bubbles can be seen in the gills, eyes and under the skin and the fish show bizarre nervous behavior.
• g. cap — a cap of gas above fluid or solid contents in a hollow viscus, e.g. in a static rumen. Seen radiologically in distended intestinal loops in paralytical ileus.
• g. edema disease — see blue wing disease.
• g. exchange — gases move by simple diffusion in response to pressure differences; net diffusion occurs from areas of high pressure to areas of lower pressure irrespective of whether the gas is present as a gas or in solution or gases moving from gas to solution or vice versa. The rate of exchange of gases in body tissues, e.g. between alveolar space and erythrocyte, is influenced by many other factors, especially the diffusion distance and the solubility of the gas.
• g. inhalation — irritant gases, e.g. manure gas, cause pulmonary edema.
• laughing g. — nitrous oxide.
• manure g. poisoning — see manure pit gas poisoning.
• tear g. — a gas that produces severe lacrimation by irritating the conjunctivae. See lacrimator.
• g. transport — relates to the efficiency of transport of gas, e.g. oxygen, by the patient as a whole. The efficiency of gas transport varies widely between normal individuals and between species, e.g. athletic breeds of horses and dogs have much faster gas transport systems than human athletes; the efficiency of gas transport in the individual depends largely on the rapidity of increase in minute ventilation, plus a similar rate of increase in cardiac output.
• g. tube — see crookes' tube.

Any aeriform or completely elastic fluid which is not a solid or a liquid. Gasses are produced by heating a liquid beyond its boiling point.

This page is about the physical properties of gas as a state of matter. For the uses of gases, and other meanings, see Gas (disambiguation).

Gas phase particles (atoms, molecules, or ions) move around freely in the absence of an applied electric field.

As a noun in the English language, a gas is one of three classical states of matter.^[1] Near absolute zero, a substance exists as a solid. As heat is added to this substance it melts into a liquid at its melting point (see phase change), boils into a gas at its boiling point, and if heated high enough would enter a plasma state in which the electrons are so energized that they leave their parent atoms from within the gas. A pure gas may be made up of individual atoms (e.g. a noble gas or atomic gas like neon), elemental molecules made from one type of atom (e.g. oxygen), or compound molecules made from a variety of atoms (e.g. carbon dioxide). A gas mixture would contain a variety of pure gases much like the air. What distinguishes a gas from liquids and solids is the vast separation of the individual gas particles as seen in the model to the right. This separation usually makes a colorless gas invisible to the human observer. The remainder of the article focuses on the interaction of gas particles in the presence of electric and gravitational fields that are considered negligible.

This article provides background information on the gaseous state of matter found between the liquid and plasma states^[2], the latter of which provides the upper temperature boundary for gases. Bounding the lower end of the temperature scale lie degenerative quantum gases^[3] which are gaining increased attention these days.^[4] High-density atomic gases super cooled to incredibly low temperatures are classified by their statistical behavior as either a Bose gas or a Fermi gas. For a comprehensive listing of these exotic states of matter see list of states of matter.

The journey begins with a brief review of the physical characteristics of gases. From that introduction, the article diverges and views gases from two distinct perspectives, macro-, or the system point of view, and microscopic, or particle viewpoint. Simplified models or gas laws are introduced next. These models characterize ideal gas behavior. From modeling, the article explores the historical connection to gases including prominent scientists along the path. The final section introduces gas terms which link to more detailed applications involving gases. Among these advanced applications are energy considerations like thermodynamics, and movement through a gas, gas flows, or aerodynamics.

Contents 1 Physical characteristics 2 Macroscopic 2.1 Pressure 2.2 Temperature 2.3 Specific volume 2.4 Density 3 Microscopic 3.1 Kinetic theory 3.2 Brownian motion 3.3 Intermolecular forces 4 Simplified models 4.1 Perfect gas 4.1.1 Calorically perfect 4.1.2 Thermally perfect 4.2 Ideal gas 4.3 Real gas 5 Historical synthesis 5.1 Boyle's Law 5.2 Law of volumes 5.3 Avogadro's Law 5.4 Dalton's Law 6 Special topics 6.1 Compressibility 6.2 Reynolds Number 6.3 Viscosity 6.4 Turbulence 6.5 Boundary layer 6.6 Maximum entropy principle 6.7 Thermodynamic equilibrium 7 Etymology 8 See also 9 Notes 10 References 11 Further reading

Physical characteristics

Drifting smoke particles provide clues to the movement of the surrounding gas.

As most gases are difficult to observe directly with our senses, they are described through the use of four physical properties or macroscopic characteristics: the gas’s pressure, volume, number of particles (chemists group them by moles), and temperature. These four characteristics were repeatedly observed by men such as Robert Boyle, Jacques Charles, John Dalton, Joseph Gay-Lussac and Amedeo Avogadro for a variety of gases in a great many settings. Their detailed studies ultimately led to a mathematical relationship among these properties expressed by the ideal gas law (see simplified models section below).

Gas particles are widely separated from one another, and as such do not influence adjacent particles to the same degree as liquids or solids. This influence (intermolecular forces) results from the magnetic charges that these gas particles carry. Like charges repel, while oppositely charged particles attract one another. Gases made from ions carry permanent charges, as do compounds with their polar covalent bonds. These polar covalent bonds produce permanent charge concentrations within the molecule while the compound's net charge remains neutral. Transient charges exist in covalent bonds of molecules and are referred to as van der Waals forces. The interaction of these intermolecular forces varies within a substance which determines many of the physical properties unique to each gas.^[5][6] A quick comparison of boiling points for compounds formed by ionic and covalent bonds leads us to this conclusion^[7]. The image to the right provides some insight into low pressure gas behavior.

Compared to the other states of matter, gases have an incredibly low density and viscosity. Pressure and temperature influence the particles within a certain volume. This variation in particle separation and speed is referred to as compressibility. This particle separation and size influences optical properties of gases as can be found in the following list of refractive indices. Finally, gas particles spread apart or diffuse in order to homogeneously distribute themselves throughout any container.

Macroscopic

Shuttle imagery of re-entry phase.

When observing a gas, it is typical to specify a frame of reference or length scale. A larger length scale corresponds to a macroscopic or global point of view of the gas. This region (referred to as a volume) must be sufficient in size to contain a large sampling of gas particles. The resulting statistical analysis of this sample size produces the "average" behavior (i.e. velocity, temperature or pressure) of all the gas particles within the region. By way of contrast, a smaller length scale corresponds to a microscopic or particle point of view.

From this global vantage point, the gas characteristics measured are either in terms of the gas particles themselves (velocity, pressure, or temperature) or their surroundings (volume). By way of example, Robert Boyle studied pneumatic chemistry for a small portion of his career. One of his experiments related the macroscopic properties of pressure and volume of a gas. His experiment used a J-tube manometer which looks like a test tube in the shape of the letter J. Boyle trapped an inert gas in the closed end of the test tube with a column of mercury, thereby locking the number of particles and temperature. He observed that when the pressure was increased on the gas, by adding more mercury to the column, the trapped gas volume decreased. Mathematicians describe this situation as an inverse relationship. Furthermore, when Boyle multiplied the pressure and volume of each observation, the product (math) was always the same, a constant. This relationship held true for every gas that Boyle observed leading to the law, (PV=k), named to honor his work in this field of study.

There are many math tools to choose from when analyzing gas properties. As gases are subjected to extreme conditions, the math tools become a bit more complex, from the Euler equations (inviscid flow) to the Navier-Stokes equations^[8] that fully account for viscous effects. These equations are tailored to meet the unique conditions of the gas system in question. Boyle's lab equipment allowed the use of algebra to obtain his analytical results. His results were possible because he was studying gases in relatively low pressure situations where they behaved in an "ideal" manner. These ideal relationships enable safety calculations for a variety of flight conditions on the materials in use. The high technology equipment in use today was designed to help us safely explore the more exotic operating environments where the gases no longer behave in an "ideal" manner. This advanced math, to include statistics and multivariable calculus, makes possible the solution to such complex dynamic situations as space vehicle reentry. One such example might be the analysis of the photo above and to the right to ensure the material properties under this loading condition are not exceeded. It is safe to say that in this flight regime, the gas is no longer behaving ideally.

Pressure

Pressurized gases.

Main article: Pressure

The symbol used to represent pressure in equations is "p" or "P" with SI units of pascals.

When describing a container of gas, the term pressure (or absolute pressure) refers to the average force the gas exerts on the surface area of the container. Within this volume, it is sometimes easier to visualize the gas particles moving in straight lines until they collide with the container (see diagram at top of the article). The force imparted by a gas particle into the container during this collision is the change in momentum of the particle. As a reminder from classical mechanics, momentum, by definition, is the product of mass and velocity.^[9] Notice that during a collision only the normal component of velocity changes. A particle traveling parallel to the wall never changes its momentum. So the average force on a surface must be the average change in linear momentum from all of these gas particle collisions. To be more precise, pressure is the sum of all the normal components of force exerted by the particles impacting the walls of the container divided by the surface area of the wall. The image to the right depicts a gas pressure and temperature spike used in the entertainment industry.

Temperature

Air balloon shrinks after submerging into liquid nitrogen

Main article: Thermodynamic temperature

The symbol used to represent temperature in equations is T with SI units of kelvins.

The speed of a gas particle is proportional to its absolute temperature. The volume of the balloon in the image to the right shrinks when the trapped gas particles slow down with the addition of extremely cold nitrogen. The temperature of any physical system is related to the motions of the particles (molecules and atoms) which make up the [gas] system.^[10] In statistical mechanics, temperature is the measure of the average kinetic energy stored in a particle. The methods of storing this energy are dictated by the degrees of freedom of the particle itself (energy modes). Kinetic energy added (endothermic process) to gas particles by way of collisions produces linear, rotational, and vibrational motion as well. By contrast, a molecule in a solid can only increase its vibration modes with the addition of heat as the lattice crystal structure prevents both linear and rotational motions. These heated gas molecules have a greater speed range which constantly varies due to constant collisions with other particles. The speed range can be described by the Maxwell-Boltzmann distribution. Use of this distribution implies ideal gases near thermodynamic equilibrium for the system of particles being considered.

Specific volume

Expanding gases link to changes in specific volume.

Main article: Specific volume

The symbol used to represent specific volume in equations is "v" with SI units of cubic meters per kilogram.

See also: Gas volume

The symbol used to represent volume in equations is "V" with SI units of cubic meters.

When performing a thermodynamic analysis, it is typical to speak of intensive and extensive properties. Properties which depend on the amount of gas (either by mass or volume) are called extensive properties, while properties that do not depend on the amount of gas are called intensive properties. Specific volume is an example of an intensive property because it is the ratio of volume occupied by a unit of mass of a gas that is identical throughout a system at equilibrium.^[11] 1000 atoms of protactinium as a gas occupy the same space as any other 1000 atoms for any given temperature and pressure. This concept is easier to visualize for solids such as iron which are incompressible compared to gases. When the seat ejection is initiated in the image above the specific volume increases with the expanding gases, while mass is conserved. Since a gas fills any container in which it is placed, volume is an extensive property.

Density

Main article: Density

The symbol used to represent density in equations is ρ (pronounced rho) with SI units of kilograms per cubic meter. This term is the reciprocal of specific volume.

Since gas molecules can move freely within a container, their mass is normally characterized by density. Density is the mass per volume of a substance or simply, the inverse of specific volume. For gases, the density can vary over a wide range because the particles are free to move closer together when constrained by pressure or volume or both. This variation of density is referred to as compressibility. Like pressure and temperature, density is a state variable of a gas and the change in density during any process is governed by the laws of thermodynamics. For a static gas, the density is the same throughout the entire container. Density is therefore a scalar quantity; it is a simple physical quantity that has a magnitude but no direction associated with it. It can be shown by kinetic theory that the density is inversely proportional to the size of the container in which a fixed mass of gas is confined. In this case of a fixed mass, the density decreases as the volume increases.

Microscopic

If one could observe a gas under a powerful microscope, one would see a collection of particles (molecules, atoms, ions, electrons, etc.) without any definite shape or volume that are in more or less random motion. These neutral gas particles only change direction when they collide with another particle or the sides of the container. By stipulating that these collisions are perfectly elastic, this substance is transformed from a real to an ideal gas. This particle or microscopic view of a gas is described by the Kinetic-molecular theory. All of the assumptions behind this theory can be found in the postulates section of Kinetic Theory.

Kinetic theory

Main article: Kinetic theory

Kinetic theory provides insight into the macroscopic properties of gases by considering their molecular composition and motion. Starting with the definitions of momentum and kinetic energy^[12], one can use the conservation of momentum and geometric relationships of a cube to relate macro system properties of temperature and pressure to the microscopic property of kinetic energy per molecule. The theory provides averaged values for these two properties.

The theory also explains how the gas system responds to change. For example, as a gas is heated from absolute zero, when it is (in theory) perfectly still, its internal energy (temperature) is increased. As a gas is heated, the particles speed up and its temperature rise. This results in greater numbers of collisions with the container sides each second due to the higher particle speeds associated with elevated temperatures. As the number of collisions (per unit time) increase on the surface area of the container, the pressure increases in a proportional manner.

Brownian motion

Random motion of gas particles results in diffusion.

Main article: Brownian motion

Brownian motion is the mathematical model used to describe the random movement of particles suspended in a fluid. The animation to the right, when in motion, illustrates how this behavior results in the spreading out of gases (entropy). These events are also described by particle theory.

Since it is at the limit of (or beyond) current technology to observe individual gas particles (atoms or molecules), only theoretical calculations give suggestions as to how they move, but their motion is different from Brownian Motion. The reason is that Brownian Motion involves a smooth drag due to the frictional force of many gas molecules, punctuated by violent collisions of an individual (or several) gas molecule(s) with the particle. The particle (generally consisting of millions or billions of atoms) thus moves in a jagged course, yet not so jagged as would be expected if an individual gas molecule was examined.

Intermolecular forces

When gases are compressed, intermolecular forces like those shown here start to play a more active role.

Main articles: van der Waals force and Intermolecular force

As discussed earlier, momentary attractions (or repulsions) between particles have an effect on gas dynamics. In physical chemistry, the name given to these intermolecular forces is van der Waals force. These forces play a key role in determining physical properties of a gas such as viscosity and flow rate (see physical characteristics section). Ignoring these forces in certain conditions (see Kinetic-molecular theory) allows a real gas to be treated like an ideal gas. This assumption allows the use of ideal gas laws which greatly simplifies the path to a solution.

Proper use of these gas relationships requires us to take one more look at the Kinetic-molecular theory (KMT). When these gas particles possess a magnetic charge or Intermolecular force they gradually influence one another as the spacing between them is reduced (model to the left illustrates one example). In the absence of any charge, at some point when the spacing between gas particles is greatly reduced they can no longer avoid collisions between themselves at normal gas temperatures found in a lab. Another case for increased collisions among gas particles would include a fixed volume of gas, which upon heating would contain very fast particles. What this means to us is that these ideal equations provide reasonable results except for extremely high pressure [compressible] or high temperature [ionized] conditions. Notice that all of these excepted conditions allow energy transfer to take place within the gas system. The absence of these internal transfers is what is referred to as ideal conditions (perfect - or well behaved) in which the energy exchange occurs only at the boundaries of the system. Real gases experience some of these collisions and intermolecular forces. When these collisions are statistically negligible [incompressible], results from these ideal equations are still valid. At the other end of the spectrum, when the gas particles are compressed into close proximity they behave more like a liquid, and hence another connection to fluid dynamics.

Simplified models

Main article: Equation of state

An equation of state (for gases) is a mathematical model used to roughly describe or predict the state properties of a gas. At present, there is no single equation of state that accurately predicts the properties of all gases under all conditions. Therefore, a number of much more accurate equations of state have been developed for gases in specific temperature and pressure ranges. The "gas models" that are most widely discussed are "Perfect Gas", "Ideal Gas" and "Real Gas". Each of these models has its own set of assumptions to facilitate the analysis of a given thermodynamic system^[13]. Each successive model expands the temperature range of coverage to which it applies. The image below and to the right illustrates one example on the successful application of these relationships in 1903. More recent examples include the 2009 maiden flights of the first solar powered aircraft, the Solar Impulse, and the first commercial airliner to be constructed primarily from composite materials, the Dreamliner.

First flight at Kitty Hawk, NC.

Perfect gas

Main article: Perfect gas

By definition, a perfect gas is one in which intermolecular forces are negligible due to the separation of the molecules and any particle collisions are elastic.

Perfect gas equation of state

The symbol n represents the number of particles grouped by moles of a substance. All other symbols in these equations use notation described earlier in the Macroscopic Section. These relationships are valid only when used with absolute temperatures and pressures.

• Chemist's version – PV = nRT

The gas constant, R, in this expression has different units than the Gas Dynamicist's version. The Chemist's version emphasizes numbers of particles (n), while the latter emphasizes the particle mass in the density term ρ.

• Gas Dynamicist's version- P = ρRT

There are two subclassifications to a perfect gas although various textbooks either omit or combine the following simplifications into a general "perfect gas" definition. For sake of clarity, these simplifications are defined separately in the following two subsections.

Calorically perfect

Main article: Calorically perfect gas

The Calorically perfect gas model is the most restrictive from a temperature perspective^[14], as it adds the following condition:

• Constant specific heats (valid for most gases below 1000 K)

u = C_vT, h = C_pT

Here u represents internal energy, h represents enthalpy, and the C terms represent the specific heat capacity at either constant volume or constant pressure, respectively.

Although this may be the most restrictive model from a temperature perspective, it is accurate enough to make reasonable predictions within the limits specified. A comparison of calculations for one compression stage of an axial compressor (one with variable C_p, and one with constant C_p) produces a deviation small enough to support this approach. As it turns out, other factors come into play and dominate during this compression cycle. These other effects would have a greater impact on the final calculated result than whether or not C_p was held constant. (examples of these real gas effects include compressor tip-clearance, separation, and boundary layer/frictional losses, etc.)

Thermally perfect

Main article: Thermally perfect gas

A thermally perfect gas is:

• in thermodynamic equilibrium
• not chemically reacting (chemical equilibrium)
• c_p - c_v = R (still valid even though specific heats vary with temperature)
• Internal energy, enthalpy, and specific heats are functions of temperature only.

u = u(T), h = h(T), du = C_vdT, dh = C_pdT

This type of approximation is useful for modeling, for example, a turbine where temperature fluctuations are usually not large enough to cause any significant deviations from the thermally perfect gas model. Heat capacity is still allowed to vary, though only with temperature and the molecules are not permitted to dissociate.^[15]

Ideal gas

An "ideal gas" is a simplified "real gas" with the assumption that the compressibility factor Z is set to 1 meaning that this pneumatic ratio remains constant. A compressibility factor of one also requires the four state variables to follow the ideal gas law.

This approximation is more suitable for applications in engineering although simpler models can be used to produce a "ball-park" range as to where the real solution should lie. An example where the "ideal gas approximation" would be suitable would be inside a combustion chamber of a jet engine^[16]. It may also be useful to keep the elementary reactions and chemical dissociations for calculating emissions.

Real gas

21 April 1990 eruption of Mount Redoubt, Alaska, illustrating real gases not in thermodynamic equilibrium.

Main article: Real gas

Each one of the assumptions listed below adds to the complexity of the problem's solution. As the density of a gas increases with pressure rises, the intermolecular forces play a more substantial role in gas behavior which results in the ideal gas law no longer providing "reasonable" results. At the upper end of the engine temperature ranges (e.g. combustor sections - 1300 K), the complex fuel particles absorb internal energy by means of rotations and vibrations that cause their specific heats to vary from those of diatomic molecules and noble gases. At more than double that temperature, electronic excitation and dissociation of the gas particles begins to occur causing the pressure to adjust to a greater number of particles (transition from gas to plasma).^[17] Finally, all of the thermodynamic processes were presumed to describe uniform gases whose velocities varied according to a fixed distribution. Using a non-equilibrium situation implies the flow field must be characterized in some manner to enable a solution. One of the first attempts to expand the boundaries of the ideal gas law was to include coverage for different thermodynamic processes by adjusting the equation to read pVⁿ = constant and then varying the n through different values such as the specific heat ratio, γ.

Real gas effects include those adjustments made to account for a greater range of gas behavior:

• Compressibility effects (Z allowed to vary from 1.0)
• Variable heat capacity (specific heats vary with temperature)
• Van der Waals forces (related to compressibility, can substitute other equations of state)
• Non-equilibrium thermodynamic effects
• Issues with molecular dissociation and elementary reactions with variable composition.

For most applications, such a detailed analysis is excessive. Examples where "Real Gas effects" would have a significant impact would be on the Space Shuttle re-entry where extremely high temperatures and pressures are present or the gases produced during geological events as in the mountain image above and to the right.

Historical synthesis

Boyle's Law

Boyle's equipment.

Main article: Boyle's Law

Boyle's Law was perhaps the first expression of an equation of state. In 1662 Robert Boyle, an Irishman, performed a series of experiments employing a J-shaped glass tube, which was sealed on one end. Mercury was added to the tube, trapping a fixed quantity of air in the short, sealed end of the tube. Then the volume of gas was carefully measured as additional mercury was added to the tube. The pressure of the gas could be determined by the difference between the mercury level in the short end of the tube and that in the long, open end. Through these experiments, Boyle noted that the gas volume varied inversely with the pressure.^[18] The image to the right shows some of the equipment Boyle used during his study of gases.

• Boyle's Law - describes a gas in which the number of particles and Temperature are constant.
• PV = constant in this situation constant = nRT from the ideal gas law.

Law of volumes

Main articles: Charles Law and Gay-Lussac's Law

In 1787, the French physicist and balloon pioneer, Jacques Charles, found that oxygen, nitrogen, hydrogen, carbon dioxide, and air expand to the same extent over the same 80 kelvin interval.

In 1802, Joseph Louis Gay-Lussac published results of similar, though more extensive experiments, indicating a linear relationship between volume and temperature. Gay-Lussac credited Charle's earlier work by naming the law in his honor. In the absence of this linkage, Dalton could have been in contention for this honor for his previously published work on partial pressures.

• Law of Volumes - Both Charles and Gay-Lussac played a role in developing this relationship^[19]
• V/T = constant - notice that constant = nR/P from the ideal gas law.

Avogadro's Law

Dalton's notation.

Main article: Avogadro's law

In 1811, Amedeo Avogadro verified that equal volumes of pure gases contain the same number of particles. His theory was not generally accepted until 1858 when another Italian chemist Stanislao Cannizzaro was able to explain non-ideal exceptions. For his work with gases a century prior, the number that bears his name Avogadro's constant represents the number of atoms found in 12 grams of elemental carbon-12 (6.022×10²³ mol^-1). This specific number of gas particles, at standard temperature and pressure (ideal gas law) occupies 22.40 liters and is referred to as the molar volume.

• Avogadro's Law - describes a gas in a container in which the pressure and temperature are constant. The simplified form for the ideal gas law follows:
• V/n = constant – notice that constant = RT/P from the ideal gas law.

Dalton's Law

Main article: Dalton's law

In 1801, John Dalton published the Law of Partial Pressures from his work with ideal gas law relationship: The pressure of a mixture of gases is equal to the sum of the pressures of all of the constituent gases alone. Mathematically, this can be represented for n species as:

Pressure_total = Pressure₁ + Pressure₂ + ... + Pressure_n

Dalton's journal is shown to the left. Among his key journal observations upon mixing unreactive "elastic fluids" (gases) were the following.^[20]:

• Unlike liquids, heavier gases did not drift to the bottom upon mixing.
• Gas particle identity played no role in determining final pressure (they behaved as if their size was negligible).

Special topics

Compressibility

Compressibility factors for air.

Main article: Compressibility factor

Thermodynamicists use this factor (Z) to alter the ideal gas equation to account for compressibility effects of real gases. This factor represents the ratio of actual to ideal specific volumes. It is sometimes referred to as a "fudge-factor" or correction to expand the useful range of the ideal gas law for design purposes. Usually this Z value is very close to unity.

Reynolds Number

Main article: Reynolds number

In fluid mechanics, the Reynolds number is the ratio of inertial forces (v_sρ) to viscous forces (μ/L). It is one of the most important dimensionless numbers in fluid dynamics and is used, usually along with other dimensionless numbers, to provide a criterion for determining dynamic similitude. As such the Reynold's number provides the link between modeling results (design) and the full scale actual conditions. It can also be used to characterize the flow.

Viscosity

Satellite view of weather pattern in vicinity of Robinson Crusoe Islands on 15 September 1999, shows a unique turbulent cloud pattern called a "von Kármán vortex street."

Main article: Viscosity

Viscosity, a physical property, is a measure of how well adjacent molecules stick to one another. A solid can withstand a shearing force due to the strength of these sticky intermolecular forces. A fluid will continuously deform when subjected to a similar load. While a gas has a lower value of viscosity than a liquid, it is still an observable property. If gases had no viscosity, then they would not stick to the surface of a wing and form a boundary layer. From the delta wing image below, it is clear that the gas particles stick to one another (see discussion on boundary layer below).

Turbulence

Air flowing past delta wing. The shadows form as the indices of refraction change within the gas as it compresses on the leading edge of this wing.

Main article: Turbulence

In fluid dynamics, turbulence or turbulent flow is a flow regime characterized by chaotic, stochastic property changes. This includes low momentum diffusion, high momentum convection, and rapid variation of pressure and velocity in space and time. The weather image to the right illustrates just one example.

Boundary layer

Main article: Boundary layer

Particles will, in effect, "stick" to the surface of an object moving through it. This layer of particles is called the boundary layer. At the surface of the object, it is essentially static due to the friction of the surface. The object, with its boundary layer is effectively the new shape of the object that the rest of the molecules "see" as the object approaches. This boundary layer can separate from the surface, essentially creating a new surface and completely changing the flow path. The classical example of this is a stalling airfoil. The image below and to the right clearly shows the boundary layer thickening as the flow moves past the point from right to left along the leading edge.

Maximum entropy principle

Main article: Principle of maximum entropy

As the total number of degrees of freedom approaches infinity, the system will be found in the macrostate that corresponds to the highest multiplicity. In order to illustrate this principle, observe the skin temperature of a frozen metal bar. Using a thermal image of the skin temperature, note the temperature distribution on the surface. This initial observation of temperature represents a "microstate." At some future time, a second observation of the skin temperature produces a second microstate. By continuing this observation process, it is possible to produce a series of microstates that illustrate the thermal history of the bar's surface. Characterization of this historical series of microstates is possible by choosing the macrostate that successfully classifies them all into a single grouping.

Thermodynamic equilibrium

Main article: Thermodynamic equilibrium

When energy transfer ceases from a system, this condition is referred to as thermodynamic equilibrium. Usually this condition implies the system and surroundings are at the same temperature so that heat no longer transfers between them. It also implies that external forces are balanced (volume does not change), and all chemical reactions within the system are complete. The timeline varies for these events depending on the system in question. A container of ice allowed to melt at room temperature takes hours, while in semiconductors the heat transfer that occurs in the device transition from an on to off state could be on the order of a few nanoseconds.

Etymology

Look up gas in Wiktionary, the free dictionary.

The word "gas" was invented by Jan Baptist van Helmont, perhaps as a Dutch pronunciation re-spelling of "chaos".^[21]

See also

• Air conditioning
• Atmosphere
• Bellows
• Breath
• Clouds
• Cooking
• Diaphragm
• Electrolysis
• Flame tests
• Fumes
• Kites
• Lighting
• Lightning
• Liquid
• Lung
• Mixmaster dynamics
• Odor
• Olfaction
• Parachute
• Petroleum
• Plasma (physics)
• Sailing
• Soaring
• Solid
• Troposhere
• Turbines
• Vapor
• Volcanic gas
• Weather
• Wind
• Wind mill
• Wind turbine

Notes

1. ^ Considering when this text was published, it thoroughly covers the three phases of matter "known" at the time. McPherson, pp.104-10
2. ^ This early 1900's discussion infers what is regarded as the plasma state. See page 137 of American Chemical Society, Faraday Society, Chemical Society (Great Britain)'s The Journal of physical chemistry, Volume 11 (Cornell - 1907).
3. ^ The work by T. Zelevinski provides another link to latest research about Strontium in this new field of study. See Tanya Zelevinsky (2009). "84Sr—just right for forming a Bose-Einstein condensate". Physics 2: 94. http://physics.aps.org/articles/v2/94. 
4. ^ ScienceDaily 4 November 2009 provides a link to more material on the Bose-Einstein condensate.
5. ^ The authors make the connection between molecular forces of metals and their corresponding physical properties. By extension, this concept would apply to gases as well, though not universally. (Cornell - 1907) pp. 164-5.
6. ^ One noticeable exception to this physical property connection is conductivity which varies depending on the state of matter (ionic compounds in water) as described by Michael Faraday in the 1833 when he noted that ice does not conduct a current. See page 45 of John Tyndall's Faraday as a Discoverer (1868).
7. ^ John S. Hutchinson (2008). Concept Development Studies in Chemistry. p. 67. http://cnx.org/content/col10264/latest/. 
8. ^ Anderson, p.501
9. ^ J. Clerk Maxwell (1904). Theory of Heat. Mineola: Dover Publications. pp. 319–20. ISBN 0486417352. 
10. ^ See pages 137-8 of Society (Cornell - 1907).
11. ^ Kenneth Wark (1977). Thermodynamics (3 ed.). McGraw-Hill. p. 12. ISBN 0-07-068280-1. 
12. ^ For assumptions of Kinetic Theory see McPherson, pp.60-61
13. ^ Anderson, pp.289-291
14. ^ Anderson, p.291
15. ^ Implies temperature limited to 1500 K. See John p.256
16. ^ John p.205
17. ^ John pp.247-56
18. ^ McPherson, pp.52-55
19. ^ McPherson, pp.55-60
20. ^ John P. Millington (1906). John Dalton. pp. 72, 77–78. 
21. ^ Online Etymology Dictionary

References

• John D. Anderson (1984). Fundamentals of Aerodynamics. McGraw-Hill Higher Education. ISBN 0070016569. 
• James John (1984). Gas Dynamics. Allyn and Bacon. ISBN 0-205-08014-6. 
• William McPherson and William Henderson (1917). An Elementary study of chemistry. 

Further reading

• Philip Hill and Carl Peterson. Mechanics and Thermodynamics of Propulsion: Second Edition Addison-Wesley, 1992. ISBN 0-201-14659-2
• National Aeronautics and Space Administration (NASA). Animated Gas Lab. Accessed February, 2008.
• Georgia State University. HyperPhysics. Accessed February, 2008.
• Antony Lewis WordWeb. Accessed February, 2008.
• Northwestern Michigan College The Gaseous State. Accessed February, 2008.

This entry is from Wikipedia, the leading user-contributed encyclopedia. It may not have been reviewed by professional editors (see full disclaimer)

Dansk (Danish)
n. - gas, luftart, benzin, speeder
v. tr. - give gas, gasforgifte, gasse
v. intr. - fylde benzintanken op, give ondt af sig

idioms:

• gas burner    gasbrænder
• gas chamber    gaskammer
• gas guzzler    benzinsluger
• gas meter    gasmåler
• gas plasma display    plasmaskærm, plasma display
• gas range    gaskomfur
• gas ring    gasapparat
• gas station    tankstation
• gas up    give gas, gasse op
• run out of gas    løbe tør for benzin

Nederlands (Dutch)
gas, benzine, tè gek iets/iemand, darmgas, lachgas, zenuwgas, opschepperij, vergassen, gas afscheiden, benzine tanken, lachgas toedienen, verblijden, opscheppen

Français (French)
n. - gaz, (Chim) gaz, (Dent) anesthésie, (Mil) gaz, (US) essence, accélérateur, (GB) bavardage (arch), rigolade, marrant (qn)
v. tr. - (gén, Mil) gazer, (GB) papoter (arch)
v. intr. - dégager des gaz, (GB) papoter, parler sans arrêt

idioms:

• gas burner    brûleur à gaz
• gas chamber    chambre à gaz
• gas guzzler    (US) voiture qui consomme beaucoup d'essence
• gas meter    compteur à gaz
• gas plasma display    (Comput) écran à plasma
• gas range    fourneau à gaz
• gas ring    (GB) brûleur à gaz, réchaud à gaz
• gas station    (US) station-service
• gas up    (US) prendre de l'essence
• run out of gas    tomber en panne d'essence, ne plus avoir d'énergie/d'idées

Deutsch (German)
n. - Gas, Benzin, Darmgase, (ugs.) leeres Gerede
v. - vergasen, auftanken, schwafeln

idioms:

• gas burner    Gasbrenner
• gas chamber    Gaskammer
• gas guzzler    Auto mit hohem Benzinverbrauch
• gas meter    Gaszähler
• gas plasma display    Plasmabildschirm, Gasplasmabildschirm
• gas range    Gasherd
• gas ring    Gasbrenner
• gas station    Tankstelle
• gas up    tanken
• run out of gas    an Kraft verlieren

Ελληνική (Greek)
n. - αέριο, φωταέριο (κν. γκάζι), αναισθητικό, (μτφ.) αερολογίες, (καθομ.) ευχάριστο ή διασκεδαστικό άτομο ή πράγμα, (ΗΠΑ) βενζίνη
v. - ναρκώνω ή θανατώνω με αέριο, φλυαρώ

idioms:

• gas burner    μπεκ (καυστήρας) γκαζιού
• gas chamber    θάλαμος αερίων
• gas guzzler    βενζινορουφήχτρα
• gas meter    ρολόι (μετρητής) γκαζιού
• gas plasma display    (Η/Υ) οθόνη πλάσματος αερίου
• gas range    κουζίνα γκαζιού
• gas ring    εστία/μάτι γκαζιού
• gas station    (ΗΠΑ) πρατήριο βενζίνης (κν. βενζινάδικο)
• gas up    γεμίζω το ντεπόζιτο
• run out of gas    μένω από βενζίνη, παύω να προκαλώ το ενδιαφέρον

Italiano (Italian)
gas, benzina, flatulenza

idioms:

• gas burner    bruciatore
• gas chamber    camera a gas
• gas guzzler    che consuma molto gas
• gas meter    contatore del gas
• gas range    cucina a gas
• gas ring    fornello a gas
• gas station    rifornitore di benzina
• gas up    fare il pieno
• inert/noble gas    gas nobile/inerte
• run out of gas    essere a secco

Português (Portuguese)
n. - gás (m), gasolina (f), conversa (f) fiada
v. - asfixiar com gás

idioms:

• gas burner    bico (m) de gás
• gas chamber    câmara (f) de gás
• gas guzzler    bebedor (m) de gasolina (coloq.) (Autom.)
• gas meter    medidor (m) de gás
• gas range    fogão (m) a gás
• gas ring    fogareiro (m)
• gas station    posto (m) de gasolina
• gas up    encher o tanque com gasolina
• inert/noble gas    gás (m) nobre (Quím.)
• run out of gas    ficar sem gás (coloq.), acabar a gasolina

Русский (Russian)
газ, горючее, болтовня, сногсшибательная шутка, газовый, отравлять газом, молоть чепуху, заправлять горючим, потрясти кого-л.

idioms:

• gas burner    газовая горелка
• gas chamber    газовая камера
• gas guzzler    неэкономичный двигатель
• gas meter    газовый счетчик
• gas range    газовая плита
• gas ring    кольцевидный рассекатель пламени на газовой горелке
• gas station    бензоколонка
• gas up    заправлять газом
• inert/noble gas    инертный газ
• run out of gas    истощить запас горючего или газа

Español (Spanish)
n. - gasolina, nafta, gas, gases, gases intestinales
v. tr. - proveer de gas o gasolina, envenenar o asfixiar con gas, impregnar con gas
v. intr. - soltar o emanar gas

idioms:

• gas burner    quemador de gas, mechero
• gas chamber    cámara de gas
• gas guzzler    automóvil que tiene un alto consumo de gasolina
• gas meter    contador/medidor de gas
• gas plasma display    pantalla de gas plasma, pantalla de plasma
• gas range    estufa de cocina a gas
• gas ring    quemador de gas, hornillo de gas
• gas station    gasolinera, estación de gasolina
• gas up    cargar gasolina/nafta
• run out of gas    acabarse la gasolina, el combustible, el gas

Svenska (Swedish)
n. - gas (i allm.), gas(bränsle), (gift)gas, lustgas, gaslåga, skrävel, bensin
v. - gasa, anfalla, förse (lysa upp) med gas, tanka (amer.)

中文（简体）(Chinese (Simplified))
气体, 可燃气, 瓦斯, 煤气, 给...供应气体, 用气体处理, 给加汽油, 用毒气攻击, 放出气体, 空谈, 加油, 吹牛

idioms:

• gas burner    煤气灶, 煤气火焰
• gas chamber    毒气室, 毒气行刑室
• gas guzzler    耗油量大的车
• gas meter    煤气表
• gas plasma display    气体等离子显示器
• gas range    瓦斯炉灶
• gas ring    环型轻便瓦斯炉
• gas station    加油站
• gas up    给汽车等加油
• run out of gas    用完的汽油, 累了

中文（繁體）(Chinese (Traditional))
n. - 氣體, 可燃氣, 瓦斯, 煤氣
v. tr. - 給...供應氣體, 用氣體處理, 給加汽油, 用毒氣攻擊
v. intr. - 放出氣體, 空談, 加油, 吹牛

idioms:

• gas burner    瓦斯灶, 瓦斯火焰
• gas chamber    毒氣室, 毒氣行刑室
• gas guzzler    耗油量大的車
• gas meter    瓦斯表
• gas plasma display    氣體等離子顯示器
• gas range    瓦斯爐灶
• gas ring    環型輕便瓦斯爐
• gas station    加油站
• gas up    給汽車等加油
• run out of gas    用完的汽油, 累了

한국어 (Korean)
n. - 기체, 휘발유
v. tr. - 가스를 공급하다
v. intr. - 가스를 내다, 잡담을 하다

idioms:

• gas up    급유하다
• run out of gas    연료를 다 쓰다

日本語 (Japanese)
n. - 気体, ガス, ガソリン, むだ話, 毒ガス, 催涙ガス, すごく楽しいこと
v. - ガス中毒させる, ガソリンを入れる, 長々とむだ話をする

idioms:

• gas burner    ガスの火口, ガスストーブ
• gas chamber    ガス処刑室, ガス室
• gas guzzler    高燃費車
• gas jet    ガスバーナー
• gas mask    ガスマスク
• gas meter    ガスのメーター, ガスメーター
• gas range    ガスレンジ
• gas ring    ガスこんろ
• gas station    ガソリンスタンド, 給油所
• gas up    タンクを満たんにする
• inert/noble gas    不活性ガス
• laughing gas    笑気
• run out of gas    ガス欠になる

العربيه (Arabic)
‏(الاسم) غاز, بنزين (فعل) يسمم بالغاز, يطلق غاز‏

עברית (Hebrew)
n. - ‮גז, בנזין, חומר מאלחש, כיף, רוח, פטפוט, תענוג, דלק, דברי-הבל (מדוברת), אדם או דבר מושכים או משעשעים (מדוברת)‬
v. tr. - ‮הרעיל בגז‬
v. intr. - ‮פטפט, קשקש, דלק‬

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